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`Energy and the Ecosystem
`
`As we near the end of the twentieth century, we look out at the vast
`universe around us and try to understand it and our place in it. For the
`most part, humans have not been very good at this. We have always tended to
`be egocentric, to assign ourselves a far larger and more central role than we
`deserve, and to insist that we must have resulted from unique and very special
`forces. Some of the history of modern science is the story of an egocentric
`people clinging desperately to their self—centered viewpoint. The opposition to
`the sun—centered universe of Kepler and Galileo was just
`that. And the
`opposition to Darwin was a refusal to admit that we share a kinship with the rest
`of life on earth and that all species have a common origin. The current pattern of
`exploiting energy resources and destroying the natural environment stems from
`a naive wishful thinking that we are somehow above the laws of physics and
`biology.
`Yet the facts remain: We are a- small part of a very large universe, subject to
`the same laws of physics, chemistry, and biology that shape the rest of the
`natural world. This means we can neither create perpetual motion machines nor
`obtain inexhaustible supplies of energy merely for the asking. It also means that
`to understand how organisms operate we must look for explanations that are
`compatible with what we know of physics and chemistry. So far we see no
`reason to believe that life violates any basic physical principle; indeed, we have
`every reason to think that we can understand life by understanding how those
`principles apply to the complex living structures we call organisms. For a long
`time, a school of thought called vitalism denied this approach. Vitalists believed
`that life somehow transcends the principles of physics and chemistry—either
`that it violates some of them or that it contains a mysterious ”something more,”
`such as a ”vital force” that accounts for organisms operating as they do. During
`the early days of modern chemistry, some scientists insisted that there must be
`”vital” processes at work in organisms that are quite different from the purely
`chemical processes seen in other systems. In fact, it was a great triumph of
`chemistry when, in 182.8, Friedrich Wbhler synthesized the organic compound
`urea from inorganic material, ammonium cyanate,
`thus showing that no
`mysterious gap separates the two realms of matter. With the growth of modern
`biochemistry, and particularly with the insights into biology obtained since
`about 1950, vitalism lost most of its force and appeal. An open—minded person
`must admit that special forces might eventually be needed to explain some
`biological phenomena, but so far that need has not arisen. A guiding principle of
`
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`CHAPTER 5
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`ENERGY AND THE ECOSYSTEM
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`' science, the principle of parsimony, tells us not to look for complex and obscure
`i explanations if simpler ones are satisfactory, and so far biology is fully
`i comprehensible with biological principles that are wholly compatible with the
`r principles of physics and chemistry.
`In this chapter, we will address some of these fundamental physical and
`t chemical concepts in relation to biology. Our chief objective is to understand
`i, energy and how it is exchanged during chemical processes, especially during
`7, metabolism. And we will examine the problems of energy exchange in a much
`broader context by taking a vast overview of the ecological relationships between
`organisms as they coexist in the environment. We begin by describing how
`organisms interact in assemblages called communities and ecosystems in which they
`get energy both from the sun and—by eating each other—from other organisms.
`Then we will come down to an atomic and molecular level, examining the
`processes by which new molecules are formed through chemical reactions. We
`will see that when a mouse eats a seed or when a fox eats a mouse, those events
`have important consequences for the populations of seed plants, mice, and
`foxes; for the seed and the mouse are soon reduced to their monomeric
`constituents, and the act of eating makes greatest sense when we understand
`how it produces the energy and chemical building blocks organisms need to stay
`alive.
`
`5-1 All organisms live in communities.
`Among the most basic features of the biological world is that nothing lives in
`i isolation. Organisms of the same kind live together in populations which are
`always interacting with populations of other kinds. They share a common living
`space and they constitute food for one another.
`Look at the familiar life in a forest. Seen from a distance, the forest is a
`patchwork of many shades of green, but standing beneath the green canopy we
`see that each patch is a different kind of tree: beech, birch, maple, ash, elm, and
`others all reach together toward the sun. And each tree is itself a little world of
`other organisms. Here is a nest of warblers, small birds with yellow, black, and
`white coloring that carefully search out the juicy caterpillars and other insects
`hidden among the leaves and bark. There is a nest of flycatchers. One of these
`solemn gray birds sits patiently, motionless on a bare twig. Suddenly it darts out
`and hangs in the sky for an instant before snatching a nearly invisible bug from
`the air and quickly returning to its perch to wait for another. Farther down one
`of the elms, a pair of woodpeckers have made their home in a hole; from there
`they fly off to patrol each tree, keeping it healthy by removing insects from its
`bark. And in a nearby maple lives a pair of hawks who occasionally add one of
`these warblers or flycatchers or woodpeckers to their diet of small mammals.
`The ground below lies thick with bushes and small plants crowned by
`bright flowers and creeping plants that spread low over the soil by stretching out
`their runners. In a clearing a patch of yellow flowers is abuzz with swarming
`bees, each stealing nectar as it unwittingly carries pollen from one flower head to
`the next, ensuring that there will be another source of nectar next year. Between
`branches at the edge of the clearing the sun sparkles off silken threads woven
`into perfect polygonal webs by silent spiders who wait for some careless insect
`to fall into their traps. In the shade of a beech tree, a patch of cool moss provides
`a green carpet interrupted only by the shimmering trails of snails who carry
`their bright—banded homes upon their backs and by an occasional beetle or
`spider who scurries across it. If you dig beneath the moss just a little you can
`"uncover a world of tiny insects, which run away as the light touches them, and
`tiger—colored millipedes that curl into spirals and urge you to go away with a
`
`5-1 ALL ORGANISMS LIVE IN COMMUNITIES
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`squirt of noxious gas. A gray mouse disappears into its home in the roots of a
`tree, while on a branch overhead a chipmunk worries a nut from its covering.
`And in the cool quiet under a cover of ferns, a silent snake stealthily tastes the
`air with its tongue until a careless tree toad hops too close and is caught with a
`lunge of the scaly head.
`The woods slope off to low, moist ground where a bright lady-slipper hangs
`over a patch of pitcher plants whose water—filled funnels trap and digest
`adventurous insects. in the shallows of a pond, green frogs with periscope eyes
`wait submerged for unwary flies to wander within tongue range. A lone deer
`steps to the edge of the pond and bends to drink, startling a wood duck who
`paddles off and then turns to snatch an insect from the clear water. A skunk
`lumbers along the pond’s edge, searching for grubs, perhaps looking for the
`duck’s nest with its prize of eggs. Nearby, a black and yellow salamander
`watches the comings and goings from a patch of violets.
`Each of these animals shares its food with parasitic worms and protists in its
`intestine, while a zoo of mites,
`ticks, and little insects lies buried in the
`mammals’ fur and the birds’ feathers. Every drop of water, every bit of soil in the
`forest,
`teems with tiny cells—bacteria, yeasts, and green algae. Here is a
`decaying log, its bottom indistinguishable from the soil it lies upon, its top
`covered with moss, its sides harboring bright shelves of fungus. There a rainbow
`of mushrooms poke their heads from the soil, their bases connected below the
`forest debris by fine mazes of white filaments. Here is a pile of deer droppings
`already covered with delicate molds that are digesting it back into the soil; and at
`its edge a carrion beetle tears loose a wad, rolls it into a ball, and pushes it off to
`provide food for the eggs she is about to lay.
`There are a thousand versions of this story. If you have never sat quietly
`and watched one of them happening, you have missed as fine a drama as has
`ever unfolded on a stage or movie screen.
`None of these organisms could live as it does without the others. They eat
`one another and are eaten. They provide shelter for one another and are in turn
`sheltered. In essence, they provide stages for one another on which each acts out
`the drama of its life. And if you are attentive enough you may even hear a love
`song or witness a mating dance.
`Populations of different organisms that live together and interact in these
`ways create a community. Its structure is roughly analogous to a household. A
`household runs on money, which is provided by those who work and which is
`used by everyone in the consumption of goods and services. The way the money
`is distributed and managed is the household's economy (eco— from Greek oi kos :
`house). Similarly, the community runs on energy, which it gets from the
`surrounding physical environment. An ecosystem, therefore, consists of the
`community plus the environment; and the relationships within the ecosystem-
`the flow of energy and the interactions of various species—are described as its
`ecology. The energy of an ecosystem comes mainly from the sun; and some
`species, primarily the green plants, are producers that bring energy into the
`system by capturing sunlight. Plants serve as food for animals and the animals
`eat each other; so the animals are the chief consumers of the system, passing
`energy from one to another as they eat. The ecosystem also contains some
`decomposers, principally molds and bacteria, that create decay and reduce dead
`organisms back to simpler materials. So, all in all, the ecosystem consists of
`producers, consumers, and decomposers living in a specific physical environ-
`ment. A coral reef, a forest, and a desert are a few examples of different types of
`ecosystems.
`To make sense of all this we must understand what energy is, why there is
`energy in sunlight, and how it can be captured by plants. We must understand,
`
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`CHAPTER 5
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`ENERGY AND THE ECOSYSTEM
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`
`too, how energy exists in chemical structures and how energy is extracted from
`the structures of plants and animals when they are used for food. We turn to this
`subject next.
`
`0 All organisms live in ecosystems which consist of a physical environment that
`harbors a community of difierent species. Members of the community can be
`generally classified either as producers, consumers, or decomposers.
`
`5-2 Energy can be transformed, but not created or destroyed.
`Energy is best understood as the ability to do work. To be sure, there are many
`kinds of work and many kinds of energy, but in the scientific sense, work is said
`to occur whenever a force acts through a distance. Thus, when you push an
`object across a table, you exert a force on it for a certain distance, and the work
`done is actually equal to the amount of force exerted times the distance traveled.
`Energy stored in your muscles is what enables you to do this work.
`Two important types of energy are potential energy, which is the energy an
`object has by virtue of its position, and kinetic energy, which is the energy of an
`object in motion. The one can be transformed into the other. Imagine, for
`example, a heavy rock lying on the ground (Figure 5-1). It has no motion of its
`own and therefore no kinetic energy. The zero level of potential energy is
`generally arbitrary, so we also say that its potential energy is zero, because it is
`as low as it can get (in this place). If you carry the rock up several flights in a tall
`building, you are using some of the energy from your body to give the rock
`greater potential energy, which increases in proportion to the height above
`ground. Now suppose you drop the rock out a window. As it falls, it continually
`loses potential energy by getting closer to the zero level on the ground. But its
`potential energy does not simply disappear; it is being transformed into kinetic
`energy, for as the rock continually gains speed it gains kinetic energy. When the
`rock hits the ground and comes to a complete standstill, its kinetic and potential
`energy return to zero. But all the kinetic energy is transformed into the kinetic
`
`5-1.
`
`The ground is taken arbitrarily as
`the level of zero energy. /in object raised
`above it has more potential energy as it
`gets higher. When itfalls, potential
`energy is converted into the kinetic energy
`of motion.
`
`As the rock falls, it loses potential energy
`and acquires kinetic energy.
`
`The faster the rock falls, the more kinetic
`energy it has. The lower it falls, the less
`potential energy it has. The sum of the two
`energies is a constant.
`
`
`
`The person uses
`chemical energy
`derived from
`food to move
`muscles. Some
`of this energy is
`given to the
`rock.
`
`Rock on the
`ground has zero
`energy.
`
`As the rock is
`carried higher, it
`acquires greater
`potential energy.
`
`As the rock hits the ground, it has maximum
`kinetic energy, which is converted into
`kinetic energy of sound (motion of air) and
`heat (random motion of molecules).
`
`5-2 ENERGY CAN BE TRANSFORMED, BUT NOT CREATED OR DESTROYED
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`energy of the soil and other things it hits; into sound energy, which is just the
`kinetic energy of air waves; and especially into some heat, for the rock and the
`ground it hits will get hotter. And this last is an important point, for heat is also a
`form of energy-it is the kinetic energy of atoms and molecules moving about at
`random (Figure 5-2.). (See Box 5-1 to learn more about how molecules move
`around.)
`In this example, we have implied that the rock has the same amount of
`energy, whether it is on the edge of the window, in mid—flight, or just hitting the
`ground. That's right. Experiments designed to test the point always lead to the
`same conclusion: The total energy in a closed systern—one that cannot gain or
`lose energy——remains constant. This is a fundamental principle of physics, the
`first law of thermodynamics or the law of conservation of energy, which says that
`energy can be transformed from one state to another but it cannot be created or
`destroyed. The potential energy of the rock at the window's edge is precisely
`equal to its kinetic energy as it touches the ground, which in turn is equal to the
`total heat, sound, and other kinetic energy after the rock hits the ground.
`Conservation laws are important in science, because many principles can
`be stated best in terms of a quantity that is never created or destroyed but that
`always remains the same, at least in the right circumstances. In Chapter 4 we
`noted that chemical equations must be balanced to satisfy the law that mass is
`always conserved, although today we say that the total of mass plus energy is
`Conserved, since mass and energy are interconvertible in atomic reactions.
`Remember also that in Chapter 3 we emphasized the conservation of genes, the
`units of heredity, during reproduction.
`Another other form of energy is electromagnetic energy, which we know
`primarily as light and radio waves. Light energy can be absorbed by matter, a
`phenomenon of the greatest biological importance—in the absorption of light
`by plants, in vision, and elsewhere. We are also interested in electrical energy.
`There is potential electrical energy between two charged particles held close
`together. For instance, since a pair of electrons tend to repel each other, they can
`do work if they are allowed to move. An electric current, consisting of moving
`electrons, can do all kinds of work. Finally—and this we will explore more.
`carefully—there is chemical energy in the bonds that hold atoms together. This
`energy can be released and converted into other forms, too. For instance, the
`chemical energy in the bonds of gasoline molecules is released when the
`gasoline is burned in an engine to produce mechanical (kinetic) energy and
`some heat.
`Several units for measuring energy are used today. Biologists and bioche-
`mists have tended to stay with the older calorie (cal), the amount of energy
`required to raise the temperature of 1 gram of water 1°C. For reference, a large
`wooden match weighs about a gram and a nickel weighs 5 grams, so this isa
`very small amount of energy. It is far more useful to measure energy with a unit
`of 1,000 calories, or 1 kilocalorie (kcal). (A growing practice that we will not
`adopt here is to measure energy in the strictly metric unit of a joule. 1 cal
`4.1840 joules.)
`0 Energy—the capacity to do worlc—can be converted from one form to another, but
`the total energy in a closed system is conserved.
`
`EXERCISE 5—1
`In an electrical generating plant, coal is burned to heat water into steam. The
`steam turns a turbine in a generator, thus creating an electric current which is‘
`used to light a bulb and turn a motor. Describe all of the transformations of one
`kind of energy into another that occur along the way. In what form will all the
`energy finally wind up?
`
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`— 110°
`— 100°
`__ 900
`_ 80..
`_ 70.,
`— 60°
`
`»: — 50°
`
`5-2. Heat is the kinetic energy of
`molecules in random motion. (a)
`Molecules in a relatively cool gas have
`little energy. They agitate molecules of the
`glass thermometer relatively little.
`Mercury atoms in the thermometer stay
`relatively compacted——they do not rise high
`in the tulJe—sa we say the temperature
`is low. (l7) Molecules in a hotter gas
`have more energy. They are moving much
`faster and agitate the molecules of the
`thermometer more, so the atoms of
`mercury move around faster and the
`column of mercury rises higher. We
`therefore say that the temperature is high.
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`1.
`
`would then have a very orderly situation, with the two
`kinds of molecules neatly arranged in their own places.
`The second law of thermodynamics says that such an
`arrangement should become more disordered, and the
`maximum state of disorder would be to have both gases
`completely mixed throughout the box. And so the natu-
`ral direction of diffusion is to carry molecules from
`a place where they are more concentrated to a place where
`they are less concentrated. The molecules of each gas
`move about randomly and independently until eventu-
`ally they are uniformly distributed. You can see this
`more easily by putting a drop of ink in a jar of water and
`leaving it undisturbed. Gradually the ink molecules will
`diffuse out until the water is uniformly colored. While
`they are diffusing, they will create a gradient of color-—
`very dense around the original drop and getting much
`lighter with distance. The ink molecules diffuse down
`their concentration gradient, in the direction of lower
`concentration. It would take energy—and would also be
`a neat trick—to make them diffuse up their gradient, so
`they became more concentrated again.
`
`Box 5-1
`
`Diffusion
`
`Molecules are continually moving. We have pointed
`out that heat is the energy of their random movement,
`and the hotter an object is the faster its molecules are
`moving. They move randomly in all directions and are
`continually colliding and bouncing away. As they do this,
`they can slip past each other so a molecule initially on
`one side of a volume can gradually move to the other
`side, just as you can move through a crowd by squeezing
`into the little spaces that open up between people. This
`process in which molecules wander about is difiusion. It
`happens even in solids. The atoms in two pieces of metal
`pressed against each other can gradually become mixed.
`It is very obvious in gases, and if someone opens a bottle
`of a smelly chemical at one end of a quiet room you can
`soon detect it at the other end. Diffusion goes on contin-
`ually in liquids, and even if cells did not move actively,
`molecules in one part would quickly move into other
`parts. Diffusion is the basic way that molecules get from
`place to place in an organism.
`Suppose we were to divide a box into two compart-
`ments,
`fill one side with oxygen and the other with
`nitrogen, and then carefully remove the barrier. We
`
`
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`BEQ 1025
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`
`5-3 Chemical reactions entail changes in energy.
`Water can be made from its elements according to the equation
`
`1-12 + %O2 e I-I20
`
`We can mix a mole of hydrogen gas with half a mole of oxygen in a calorimeter
`process) and give the
`(an instrument that measures the heat released in a
`mixture a tiny spark to start the reaction. There is an explosion, the hydrogen
`and oxygen are converted to steam, and the calorimeter shows that 58 kcal of
`heat have been released.
`The law of conservation of energy must hold true here. A certain amount of
`energy, in the hydrogen and oxygen, went into the calorimeter. Some heat came
`out. Therefore the steam inside must have less energy than was contained in the
`original gas mixture. This experiment tells us, in fact, that a mole of water
`always contains 58 kcal less energy than does an equivalent amount of hydrogen
`and oxygen. The difference is called the heat of formation of water. During this
`reaction, the bonds between hydrogen atoms and the bonds between oxygen
`atoms are broken, and bonds that hold hydrogen and oxygen atoms together are
`formed. Because energy is released, the total energy in all the bonds in water
`molecules must be less than in the bonds of hydrogen and oxygen molecules;
`and by doing experiments of this kind with many different materials one can
`determine just how much energy is in each type of bond. For instance, the I-l—H
`bond in hydrogen contains 103 kcal/mole and the O—-H bond contains 109
`kcal/ mole.
`The way in which energy is exchanged during a chemical reaction can seem
`confusing unless the situation is stated precisely. When two atoms are bonded,
`
`5--3
`
`CHEMICAL REACTIONS ENTAIL CHANGES IN ENERGY
`
`12.3
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`they are in a stable molecular arrangement. Therefore, energy must be added to
`the molecule to break the bond, and this, by definition, is the amount of energy
`”in" the bond. When a new bond is made, energy is released as the atoms come
`together in a new, stable arrangement. Now we can see more clearly what is
`happening when water forms from hydrogen and oxygen (Figure 5-3). It takes
`103 kcal to break the bonds in the mole of hydrogen and% X 117, or 58.5 kcal, to
`break the bonds in half a mole of oxygen. These molecules are reassembled into
`a mole of water, and when each of the 0-H bonds in each molecule is formed,
`Z X 109, or 218 kcal, are released. The difference is 161.5 — 218 = -56.5 kcal
`(very close to the measured value of 58 kcal), where the negative value indicates
`that energy is released. Thus, water is shown at a lower point on the energy scale
`of Figure 5-3 than the gases from which it is made.
`Our earlier example of the falling rock points out another general rule: that
`a natural tendency of matter is to go from a position of higher potential energy
`to one of lower potential energy. In nature, everything tends to run downhill,
`not uphill. Similarly, a chemical reaction tends to go from a higher to a lower
`energy state. Just as the rock is most stable when lying on the ground, the
`hydrogen and oxygen atoms are in a more stable condition in water than they
`are in the two gases, because more energy is required to break all the bonds in
`water. During the chemical transformation, energy comes out in the form of
`heat, so this is called an exothermic reaction (ex- : out, therm = heat). We expect
`exothermic reactions to occur naturally and spontaneously because they
`correspond to the natural tendency of a process to run downhill. The opposite
`
`The energy released when hydrogen
`5-3.
`and oxygen atoms form water molecules.
`
`Reactant molecules (H2 and 02)
`
`G19
`.-
`
`
`
`
`
`
`
`Energylevel(kcal/mole)
`
`H20(gas)
`
`124
`
`—58kcal/mole
`
`218 kcal
`
`~'~—
`
`CHAPTER 5
`
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`5-4. As water gets warmer, its
`molecules move about more randomly and
`the system becomes more disordered.
`
`i
`.9
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`or.
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`5-4 ALL SYSTEMS TEND TO BECOME MORE DISORDERED
`
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`f process, in which hydrogen and oxygen are liberated from water, is called an
`endothermic reaction (end- = in) and is equivalent to running energetically uphill.
`, This endothermic reaction can only happen if energy is put into the water in an
`appropriate way—for instance, by passing an electric current through it.
`Strictly speaking, when processes occur at constant pressure, as they do in
`organisms, the energy that is being exchanged is called enthalpy or heat content,
`p denoted by H. It differs very little from the energy measured by a calorimeter.
`L’ The standard Creek A (delta) is used to denote ”change in” H:
`
`AH : Hot all products - Hot all reactants
`
`Then if AH is negative the reaction is exothermic and if it is positive the reaction
`is endothermic.
`This would appear to give us a criterion for deciding whether a process will
`occur naturally and spontaneously. In such reactions, heat should be given off,
`while in those that do not occur naturally it should be necessary to add some
`outside heat to force the process to occur. And this tends to be true. Heat is given
`off in most processes that occur naturally, although some processes violate the
`rule. For instance, when ammonium chloride dissolves in water, quite sponta-
`neously, the beaker gets cold as heat is taken up out of the water. Here is an
`endothermic process that occurs naturally, apparently in violation of a law of
`nature. But since, by definition, laws of nature cannot be violated, this means
`that there must be some other laws that account for the exceptions. We explore
`one of these in the next section.
`
`0 All chemical reactions tend to create structures that have less energy——that is,
`more stable bonds—than the structures they were made from.
`
`5-4 All systems tend to become more disordered.
`Although the tendency to achieve a lower state of energy is a powerful force, an
`equally powerful but somewhat
`less obvious force is
`the tendency for
`everything to become more disordered. This is our natural experience, and
`much of our lives are spent trying to counteract this force. Thus we repeatedly
`organize our desks, arrange the clothes in our dresser drawers, and remove
`accumulations of trash and dirty dishes. Orderly piles of papers or clothes tend
`T to fall into disordered heaps, not the other way around. And so it is with atoms
`T and molecules.
`We can illustrate this nicely using water as an example. Figure 5~4 shows
`At that in an ice crystal the water molecules are highly ordered, each one held to its
`T neighbors by hydrogen bonds. As the ice melts, some molecules can break free
`_ of the crystal lattice. (Remember that heat
`is just
`the random motion of
`T molecules.) As the ice warms up, its molecules vibrate and turn faster, and some
`, of them move so fast that they can break away from the crystal because their
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`kinetic energy is greater than the energy of the hydrogen bond. Even in liquid
`water, the molecules still tend to form hydrogen-bonded ”icebergs” whose size
`and shape are always changing as some molecules leave them and others make
`new bonds. Water is clearly more disorganized than ice because many
`molecules are moving freely and randomly. Then as water gets hot enough to
`boil, many molecules acquire enough kinetic energy to fly away from the rest
`completely, forming a water vapor that is even more disorganized than liquid
`water.
`
`Now the natural tendency of ice, at temperatures above its melting point, is
`to become disorganized. Even though the energy of each molecule is increas-
`ing—and this seems to be just the opposite of the natural course of things—the
`tendency for things to become disordered is strong enough to override the
`tendency to become less energetic. To understand this drive toward disorder,
`we define a quantity called the entropy of a system, S, which is closely related to
`the degree of disorder. Because the relationship is rather complex we won't try
`to define entropy strictly; it is enough to know that a highly ordered system like
`ice has low entropy and a disordered one like water vapor has high entropy. The
`second law of thermodynamics states that in all natural processes, the entropy of an
`isolated system (one that cannot exchange matter or energy with its surround—
`ings) tends to increase.
`Whenever two independent principles are in operation, they may conflict
`with each other. On the one hand, systems tend to acquire lower energies; on the
`other, they tend to acquire higher entropies. If there were a conflict between
`these tendencies, which one would win? We just noted that water acquires more
`energy as it acquires higher entropy, so in this case entropy wins out. But why
`not the other way around?
`The problem is solved by defining another quantity, the free energy, denoted
`by G. Free energy is defined by G : H —— TS, where T is temperature (the
`absolute or Kelvin temperature scale on which 0° corresponds to ——273.16°C).
`When any process occurs at a constant temperature, the change in free energy is
`AG 2 AH — TAS. Now all the laws of thermodynamics are satisfied if G
`decreases in every process—if AG is negative. Generally, AH is negative (21
`decrease in energy) and AS is positive (an increase in disorder); since there is a
`minus sign in front of the entropy term, the result is a negative AG. However,
`there can be different types of processes (Figure 5-5). In some, AH may be
`positive but the increase in entropy is great enough to dominate and make AG
`negative. In others, AS may be negative, signaling a decrease in order, but this
`will be overridden by a large negative AH.
`In sum, the natural direction of a process is influenced by the tendency
`toward minimum energy and the tendency toward maximum disorder. The free
`energy takes both into account. In a spontaneous process, free energy will
`decrease; such a process is called exerganic. In an endergonic process, the free
`energy increases. Such processes do not occur by themselves, but they occur all
`
`—TAS§
` .—.«..m.+~..u~w!.»..‘sr.>.:»w~:'a:-<5’-'::. D I
`
`
`AG
`
`126
`
`CHAPTER 5
`
`ENERGY AND THE ECOSYSTEM
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`Possible changes in energy and
`5-5.
`entropy in difierent processes.
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`the time in organisms; and we will see in the next section that they can be made
`to occur if something else drives them.
`The free energy has one other meaning. One can show that the maximum
`. amount of useful work that can be done by any process is equal to the free
`“energy change that occurs. This is important in biology because metabolic
`rocesses are carried out in order to get energy for work, and the change in free
`nergy in a reaction tells us how much work it can perform.
`
`Chemical reactions tend to go in the direction of decreasing free energy-
`resulting in compounds with a reduced ability to do work,
`
`5-5 Endergonic processes can be driven by coupling
`0 exergonic processes.
`Now consider the entropy of a living organism. Its constituents are primarily
`olymeric organic compounds with a high degree of order. A protein, for
`instance, is a specific arrangement of amino acids, one of an enormous possible
`number of such sequences, and its constituent monomers are all highly ordered
`rrangements